NATIONAL OPEN UNIVERSITY OF NIGERIA SCHOOL OF … 391.pdf · i national open university of nigeria school of science and technology course code: chm 391 course title: practical chemistry

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NATIONAL OPEN UNIVERSITY OF NIGERIA

SCHOOL OF SCIENCE AND TECHNOLOGY

COURSE CODE: CHM 391

COURSE TITLE: PRACTICAL CHEMISTRY V - INORGANIC AND

ANALYTICAL

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CHM 391: PRACTICAL CHEMISTRY V - INORGANIC AND

ANALYTICAL

Course Writer Dr. Kelle Henrietta Ijeoma

Chemistry Unit

School of Science and Technology

National Open University of Nigeria (NOUN)

Course Editor Dr. Ogunsipe Abimbola

Dept. of Chemistry

Federal University of Petroleum Resources (FUPR)

Effurun, Warri, Delta State.

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CHM 391 COURSE GUIDE

CONTENTS PAGE

Introduction--------------------------------------------------------------- iii

Course Description ------------------------------------------------------ iii

What you will learn in this Course ------------------------------------ iii

Course Aims -------------------------------------------------------------- iv

Course Objectives -------------------------------------------------------- iv

Working through this Course ------------------------------------------- iv

Course Materials ---------------------------------------------------------- v

Study Units ---------------------------------------------------------------- v

Textbook and References ------------------------------------------------ v

Assessment ----------------------------------------------------------------- vi

Summary ------------------------------------------------------------------- vi

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INTRODUCTION

CHM 391 (practical chemistry v-inorganic and analytical) is a practical chemistry course that

exposes the students to the theory of useful analytical procedures involving quantitative and

qualitative techniques. You are familiar with the theoretical concepts of titrimetric (volumetric)

analysis, analytical techniques involving spectroscopic methods as well as the concept of errors

as it relates to analytical chemistry, and different statistical tools for data handling.

The 300- level inorganic chemistry syllabus comprises the periodic table and the compounds

formed from their elements, with the exception of those formed between carbon and hydrogen.

CHM 391 provides you the opportunity to apply the theoretical knowledge acquired in CHM 202

to quantitatively and qualitatively analyse elements of the periodic table as well as their

compounds. It will enable you acquire the sense of collecting and assimilating the clues to

determine the composition of unknown substances.

COURSE DESCRIPTION

This is basically an analytical/inorganic practical course. The analytical chemistry aspect deals

with some spectroscopic technique employed in qualitative and quantitative estimation. The

inorganic part of the practical is based on the use of classical methods of analysis to obtain

quantitative and qualitative data/information of an analyte of interest.

In both parts; analytical and inorganic practical, mainly inorganic elements and their compounds

will be determined.

WHAT YOU WILL LEARN IN THIS COURSE

In this course you will learn about the principle of some spectroscopic methods as well as

classical methods of analysis. You will be able to follow the experimental procedure outlined for

each spectroscopic and classical method mentioned in this course to obtain the quantitative and

qualitative information/data you are asked to determine. This will help you develop analytical

skills required in analyzing an analyte.

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COURSE AIMS

The course aims to give you the opportunity to explore different analytical techniques viz

spectroscopic and classical techniques available for quantitative and qualitative elucidation of the

composition of an unknown sample. In so doing, you learn to transfer your theoretical

knowledge to solving practical problems. This course also aims to help you develop analytical

skills.

COURSE OBJECTIVES

At the end of this course, you should be able to

Describe the principle of UV visible spectroscopy

Use UV visible spectroscopy to estimate the concentration of an analyte

Explain the principle of colometric analysis

Determine the concentration of a coloured substance by colometric method

Explain the principle of Infrared (IR) spectroscopy

Determine the functional groups present in a given sample

Determine the concentration of an analyte by Infrared spectroscopic technique

Explain the principle of Atomic Absorption Spectroscopy (AAS)

Use Atomic Absorption Spectroscopic technique to determine the concentrations of metal ions in solution

Explain the principle of precipitation gravimetry

Use precipitation gravimetric technique to separate an analyte from a sample solution and determine the quantity of analyte present in a sample

Explain the principle of qualitative inorganic analysis

Separate cations of a group from another group and identify the cations in a group

Explain the principle of potentiometric titration

Determine the end point of a redox titration by potentiometry

Explain acidity in water, sources and significance of acidity in water

Determine acidity in water

Explain alkalinity in water, sources and significance of alkalinity in water

Determine alkalinity in water

Explain hardness in water, sources and significance of hardness in water

Determine hardness in water

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WORKING THROUGH THIS COURSE

The course is divided into two modules which are subdivided into 10 units. It is required that you

study the units in details and carry out the experiments contained herein. An instructor will guide

you through the practical class.

COURSE MATERIALS

You will be provided with the following materials:

1. Course Guide

2. Study Units

STUDY UNITS

The following are the two modules and ten units contained in this course:

Module 1

Unit 1 UV- Visible Spectroscopy

Unit 2 Colorimetry

Unit 3 Infrared Spectroscopy

Unit 4 Atomic Absorption Spectroscopy

Module 2

Unit 1 Precipitation Gravimetry

Unit 2 Qualitative Analysis of Cations

Unit 3 Potentiometric Titration

Unit 4 Determination of Acidity in Water and Waste Water

Unit 5 Determination of Alkalinity in Water

Unit 6 Determination of Hardness in Water

Carrying out practical is very essential for understanding the theoretical concepts learned. This

helps to move your cognitive level from knowledge to comprehension, application and synthesis.

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Module 1 deals with some different spectroscopic techniques employed in analytical chemistry.

These techniques include UV visible, colorimetric, infrared (IR) and atomic absorption

spectroscopic techniques. Here the principles of the analytical techniques mentioned are

explained. A practical session precedes each principle.

Module 2 deals with classical methods of analysis. The classical methods dealt with here are;

precipitation gravimetry, qualitative inorganic analysis of cations, determination of acidity in

water, alkalinity as well as hardness in water using titrimetric method .The principle of

precipitation gravimetry, qualitative inorganic analysis of cations,, as well as acidity, alkalinity

and hardness in water is explained briefly. This is preceded by a practical session.

TEXTBOOKS AND REFERENCES

Below are some recommended textbooks you may wish to consult to enhance your learning. You

may also need to exploit other e- reading facilities such as the internet.

Mendham, J., Denney, R.C., Barnes, J.D., and Thomas, M.J.K., (2008), Vogels Textbook of

Quantitative Chemical Analysis, 6th

Edition, Pearson Education.

Gary, D.C., (1980), Analytical Chemistry, 3rd

Edition, John Wiley & Sons, New York.

Braun, R.D., (1983), Introduction to Chemical Analysis, McGraw Hill, Auckland.

Harvey, D.,(2000), Modern Analytical Chemistry, Mcgraw Hill Higher Education Companies,

Boston.

ASSESSMENT

There are two aspects of assessment for this course: the tutor-marked assignment (TMA) and the

end of course examination.

The TMAs shall constitute the continuous assessment component of the course. They will be

marked by the tutor and equally account for 30% of the total course score. Each learner shall be

examined in four TMAs before the end of course examination.

The end of course examination shall constitute 70% of the total course score.

SUMMARY

CHM 391 analytical and inorganic practical course is based on theoretical inorganic and

analytical courses. Students are expected to make use of the theoretical knowledge acquired from

these courses to determine the quantitative and qualitative composition of elements of the

periodic table and their elements exception of those formed between carbon and hydrogen.

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MAIN COURSE

CONTENT PAGE

Module 1 1

Unit 1 .. 1

Unit 2 .. 21

Unit 3 ............................................................................................................. 29

Unit 4 .. 40

Module 2 .. 48

Unit 1 . 48

Unit 2 ................................................................................................................ 68

Unit 3 . 79

Unit 4 ...... 87

Unit 5 ... 97

Unit 6 108

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MODULE ONE (ANALYTICAL)

UNIT 1 UV- VISIBLE SPECTROSCOPY

1.0 INTRODUCTION

2.0 OBJECTIVES

3.0 MAIN CONTENT

3.1 PRINCIPLE OF UV- VISIBLE SPECTROSCOPY

3.2 KINDS OF MOLECULES THAT CAN ABSORB UV- VISIBLE RADIATION

3.2.1 THE KINDS OF POSSIBLE TRANSITION AN ORGANIC MOLECULE CAN UNDERGO

3.2.2 ELECTRONIC TRANSITION INVOLVING INORGANIC COMPOUNDS

3.3 APPLICATIONS OF UV-VISIBLE SPECTROSCOPY

3.4 DETERMINATION OF THE CONCENTRATION OF AN ANALYTE USING UV- VISIBLE SPECTROSCOPY

3.4.1 USING UV-VISIBLE ABSORPTION SPECTRA TO FIND CONCENTRATION

3.4.1.1 FINDING CONCENTRATION USING MOLAR ABSORPTIVITY

3.4.1.2 DETERMINATION OF WAVELENGHT OF ABSORPTION FOR THE

PREPARATION OF CALIBRATION CURVE

3.4.1.3 FINDING CONCENTRATION BY PLOTTING A CALIBRATION

CURVE

3.5 IDENTIFICATION OF AN ANALTYTE USING UV-VISIBLE

SPECTROSCOPY

3.6 BRIEF INTRODUCTION TO A SPECTROPHOTOMETER

3.7 BRIEF DESCRIPTION OF HOW TO USE A UV-VISIBLE

SPECTROPHOTOMER

3.8 EXPERIMENTAL

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4.0 CONCLUSION

5.0 SUMMARY

6.0 TUTOR MARKED ASSIGNMENT

7.0 REFERENCES/ FURTHER READING

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1.0 INTRODUCTION

Spectroscopy is a major branch of analytical chemistry that deals with the study of concentration

of analyte as a function of amount of radiation absorbed when electromagnetic radiation from

appropriate source is directed at it. Spectroscopy is also defined as a method of analysis, which

involves the measurements of the intensity and wavelength of radiation that is either absorbed or

transmitted. What you need to understand from both definitions is that, when electromagnetic

radiation passes through a solution of a compound (sample), a certain amount of light radiation is

absorbed by the molecules. According to Beers law, the amount of radiation absorbed by the

molecule in solution is proportional to the number of absorbing molecules in the solution

(concentration). The absorbed radiation brings about a decrease in the intensity of the transmitted

(unabsorbed) radiation. The more the number of absorbing molecules (concentration), the greater

is the intensity of absorption. Spectroscopic techniques can be used in the determination of the

concentration (quantitative analysis) of an analyte and in the identification (qualitative analysis)

of an analyte. The instruments used to study or measure the absorption or emissions of

electromagnetic radiation as a function of wavelength are spectrometers and spectrophotometers.

2.0 OBJECTIVES

After studying this unit, you should be able to:

Explain the fundamental principle behind spectroscopy

Discuss/ explain UV- Visible spectroscopy

State and explain the principle of UV- Visible spectroscopy

State the uses of UV-Visible spectroscopy

Prepare standard solution

Construct calibration curve based on Beers law

Explain how UV-Visible spectroscopy can be used to determine the concentration of an analyte and identification of an analyte

Carry out experiments which involve the use of UV- Visible spectroscopy to analyse a sample

3.0 MAIN CONTENT

3.1 PRINCIPLE OF UV- VISIBLE SPECTROSCOPY

The electromagnetic radiation covers a long range of radiations which are broken down into

different regions according to wavelength. The Ultraviolet (UV) region extends from about 10 to

380 nm, but the most analytically useful region is from 200 380 nm, called the near ultraviolet

region. Below 200 nm, the air absorbs appreciably and so the instruments are operated under a

vacuum; hence, this wavelength region is called the vacuum ultraviolet region. The visible

region is the region of wavelengths that can be seen by the eye, that is, the light appears as a

colour. It extends from the near ultraviolet region (380 nm) to about 780 nm. When an organic

molecule absorbs UV- Visible radiation, the energy from UV or visible light causes the outer

electrons from a lower energy to be raised to a higher energy level, corresponding to an

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electronic transition. This transition of electrons is from molecular bonding orbital to the higher

energy antibonding molecular orbital. According to the molecular orbital theory, the shared

electron pair of a covalently bonded atoms may be thought of as occupying molecular orbitals

(MO) which is of a lower energy and has a corresponding unoccupied orbitals called antibonding

molecular orbitals, these correspond to excited state energy levels (higher energy level). When

the molecule is in the ground state, both electrons are paired in the lower-energy bonding orbital

this is the Highest Occupied Molecular Orbital (HOMO). The antibonding, in turn, is the

Lowest Unoccupied Molecular Orbital (LUMO).

When organic molecules that are capable of absorbing UV- Visible radiation are exposed to the

radiation at a wavelength with energy equal to the difference between the energy of the HOMO

and LUMO (with energy equal to E, the HOMO-LUMO energy gap) this wavelength will be

absorbed and the energy used to bump one of the electrons from the HOMO to the LUMO. The

energy absorbed appears as absorption peaks at the wavelength it corresponds to on the UV

Visible spectrum. The extent of absorption of electromagnetic radiation corresponds to the

concentration of the analyte through the application of Beer Lambert law.

Molecules containing pi electrons or non-bonding electrons can absorb the energy in the form of

ultraviolet or visible light to excite these electrons to higher antibonding molecular orbitals.

The more easily excited the electrons (i.e. lower energy gap between the HOMO and the LUMO)

the longer the wavelength of light it can absorb.

3.2 KINDS OF MOLECULES THAT CAN ABSORB UV-VISIBLE RADIATION

Not all molecules absorb UV Visible radiation. When the energy gap between the HOMO and

LUMO is large absorption will not take place, but when this energy gap is small absorption will

take place. The electronic transitions that take place in the visible and ultraviolet regions of the

electromagnetic radiation are due to absorption of radiation by specific types or groups, bonds,

and functional groups within the molecule, known as chromophores. These contain valence

electrons of low excitation energy.

3.2.1 The kinds of possible transitions an organic molecule can undergo are:

* Transitions

An electron of a saturated compound in a bonding orbital is excited to the corresponding

antibonding orbital *. The energy required is large. The high excitation energy makes it unable

to contribute to absorption in the visible or uv regions. For example, methane (which has only

C-H bonds, and can only undergo * transitions) shows an absorbance maximum at 125

nm. Absorption maxima due to * transitions are not seen in typical UV-Vis. spectra (200 -

700 nm).

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n* Transitions

Saturated compounds containing atoms with lone pairs (non-bonding electrons) are capable of

n* transitions. These transitions usually need less energy than

* transitions because

the non bonding electrons are less tightly held than sigma electrons, however, the energy gap

between the HOMO and LUMO is large. n* transitions occur at wavelengths less than

200nm.The number of organic functional groups with n* peaks in the UV region is small.

n* and

* Transitions

Unsaturated compounds containing atoms with lone pairs (non-bonding electrons) are capable

of n*

transition. The non bonding electrons are less tightly held, they can be excited by

visible or uv radiation to unoccupied pi antibonding orbital (*)

. Electrons in pi orbitals of an

unsaturated compound are responsible for * Transitions. They are the most readily excited

and responsible for a majority of electronic spectra in the visible and uv regions.

It is important to mention at this point that, most absorption spectroscopy of organic compounds

are based on transitions of n or electrons to the * excited state. This is because the absorption

peaks for these transitions fall in an experimentally convenient region of the spectrum (200 - 700

nm). The energy gap between the HOMO and LUMO is small.

Generally the compounds that absorb strongly UV-visible radiation are molecules with

conjugated pi systems (compounds having more than one double bond, alternating a single

bond). In these group, the energy gap for *

transitions is smaller than for isolated double

bonds, and thus the wavelength absorbed is longer.

Molar absorptivities from n* transitions are relatively low, and range from 10 to100 L mol

-1

cm-1

. * transitions normally give molar absorptivities between 1000 and 10,000 L mol

-1

cm-1

.

3.2.2 Electronic transition involving inorganic compounds

Charge - Transfer Absorption

Many inorganic species show charge-transfer absorption and are called charge-transfer

complexes. For a complex to demonstrate charge-transfer behaviour, one of its components must

have electron donating properties and another component must be able to accept electrons.

Absorption of radiation then involves the transfer of an electron from the donor to an orbital

associated with the acceptor.

The intense colour of metal chelates is frequently due to charge transfer transitions. This is

simply the movement of electrons from the ligand to metal ion or vice versa. Such transitions

include promotion of electrons from levels in the ligand or from bonding orbitals to the

unoccupied orbitals of the metal ion or promotion of bonded electrons to unoccupied orbitals

of the ligand. When such transitions occur, a redox reaction actually occurs between the metal

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ion and the ligand. Usually, the metal ion is reduced and the ligand is oxidized, and the

wavelength (energy) of maximum absorption is related to the ease with which the exchange

occurs.

Molar absorbtivities from charge-transfer absorption are large (greater that 10,000 L mol-1

cm-1

).

3.3 APPLICATIONS OF UV-VISIBLE SPECTROSCOPY

UV-Vis spectroscopy is routinely used in analytical chemistry for the quantitative determination of different analytes, such as highly conjugated organic compounds,

transition metal ions and biological macromolecules.

UV-Vis spectroscopy can be used to determine the concentration of an analyte in a solution. This is based on Beer Lambert law. The law is simply an application of the

observation that, within certain ranges, the absorbance of a chromophore at a given

wavelength varies in a linear fashion with its concentration: the higher the concentration

of the molecule, the greater its absorbance. If we divide the observed value of A at max

by the concentration of the sample (c, in mol/L), we obtain the molar absorptivity, or

extinction coefficient (which is a characteristic value for a given compound The wavelengths of absorption peaks can be correlated with the types of bonds in a given

molecule and are valuable in determining the functional groups within a molecule.

3.4 DETERMINATION OF THE CONCENTRATION OF AN ANALYTE USING UV-

VISIBLE SPECTROSCOPY

The absorption of light radiation by solutions can be elucidated by a combination of the

laws of Beer and Lambert. These two laws relate the absorption to concentration and to the

thickness of the absorbing layer respectively.

Beers law states that the absorption of light is directly proportional to the number of the

absorbing molecules. That is, the transmittance decreases exponentially with the number or

concentration of the absorbing molecules.

Mathematically, Beers law is represented as:

C or

c

Where,

is the absorbance (A), C is the concentration and is a constant, I0 = incident

light, I = transmitted light

This can also be represented as: A C or A = C

Lamberts law states that same proportion of incident light is absorbed per unit thickness

irrespective of its intensity, and that each successive unit layer absorbs the same proportion of

light falling on it. For example, if the incident light is 100% and 50% of it is absorbed per unit

layer, the intensity of light will decrease exponentially as follows: 50%, 25%, 12.5%, 6.25%, etc.

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Thus, according to this law,

Or

= k

Note that: absorbance (A) =

= k

Where, k is a constant and is the path length.

The two laws are combined together and called Beer-Lamberts law:

A =

or

=

Where, is a constant called molar extinction coefficient (or molar absoptivity) which is

numerically equal to the absorbance of a molar solution in a cell of 1cm path length.

Note: While Lamberts law holds for all cases, Beers law is only obeyed by dilute solutions.

Absorbance and transmittance

The absorbance (A) is the measure of the fraction of light radiation that is absorbed by a given

sample solution, while transmittance (T) is the fraction of incident light that is not absorbed

(i.e. transmitted by the solution).

Transmittance (T) =

But, absorbance is related to transmittance as follows:

A = - T =

Note that as the absorbance of a solution increases, the transmittance decreases.

3.4. 1 Using UV- Visible absorption spectra to find concentrations

A =

The absorbance (

) is measured by a spectrometer or spectrophotometer.

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3.4.1.1 Determination of concentration using the molar absorptivity

If you know the molar absorptivity of a solution at a particular wavelength, and you measure the

absorbance of the solution at that wavelength, it is easy to calculate the concentration. The only

other variable in the expression above is the length of the solution. That's easy to measure and, in

fact, the cell containing the solution may well have been manufactured with a known length of 1

cm (most cells are manufactured with a known length of 1 cm).

For example, let's suppose you have a solution in a cell of length 1 cm. You measure the

absorbance of the solution at a particular wavelength using a spectrometer. The value is 1.92.

You find a value for molar absorptivity in a table of 19400 for that wavelength.

Substituting those values:

A =

1.92 = 19400 x 1 x c

C = 1.92

19400

= 9.90 x 10-5

mol dm-3

Notice what a very low concentration can be measured provided you are working with a

substance with a very high molar absorptivity.

This method, of course, depends on you having access to an accurate value of molar absorptivity.

It also assumes that the Beer-Lambert Law works over the whole concentration range.

It is much better to measure the concentration by plotting a calibration curve.

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3.4.1.2 Determination of wavelength of absorption for the preparation of calibration curve

In preparing the calibration curve, it is imperative to know the wavelength at which the

absorbance of the different standard solutions and the sample solution is to be determined. The

wavelength is determined by plotting absorbance against wavelength (measured in nanometers).

The wavelength corresponding to the absorbance maximum (or transmittance minimum) is read

from the plot and used to prepare the calibration curve. Let us take an example. Suppose you are

asked to prepare a calibration curve to determine the concentration of a sample whose

concentration is unknown by UV- Visible spectroscopic method. What you will do to determine

the wavelength at which this determine is to be carried out is to prepare a standard solution of the

sample, after which you measure the absorbance and percentage transmittance over a series of

wavelengths e.g wavelengths covering the range 360 nm 640 nm at for instance 10nm intervals

i.e. 360 nm, 370 nm, 380 nm e.t.c. Plot the graph of absorbance against the wavelength or

percentage transmittance. You would obtain a graph as shown in Figure 1

Figure 1

The max of 524 nm is the wavelength of maximum absorption so the calibration curve will be

determined at this wavelength.

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3.4.1.3 Finding concentration by plotting a calibration curve

Doing it this way you don't have to rely on a value of molar absorptivity, the reliability of the

Beer-Lambert Law, or even know the dimensions of the cell containing the solution.

What you do is make up a number of solutions of the compound you are investigating - each of

accurately known concentration. Those concentrations should bracket the concentration you are

trying to find - some less concentrated; some more concentrated.

For each solution, you measure the absorbance at the wavelength of maximum absorption - using

the same container for each one. Then you plot a graph of absorbance against concentration. This

is a calibration curve.

According to the Beer-Lambert Law, absorbance is proportional to concentration, and so you

would expect a straight line. That is true as long as the solutions are dilute, but the Law breaks

down for solutions of higher concentration, and so you might get a curve under these

circumstances.

As long as you are working from values either side of the one you are trying to find, that is not a

problem.

Having drawn a line of best fit, the calibration curve will probably look like Figure 2, and it's

what you will probably get if you are working with really dilute solutions.

Figure 2

Notice that no attempt has been made to force the line back through the origin. If the Beer-

Lambert Law worked perfectly, it would pass through the origin, but you can't guarantee that it is

working properly at the concentrations you are using.

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Now all you have to do is to measure the absorbance of the solution (analyte in solution) with the

unknown concentration at the same wavelength as used for the preparation of the calibration

curve. If, for example, it had an absorbance of 0.600, you can just read the corresponding

concentration from the graph as below (Figure 3).

Figure 3

3.5 IDENTIFICATION OF AN ANALTYTE USING UV-VISIBLE SPECTROSCOPY

If you compared the peaks on a given UV-visible absorption spectrum with a list of known

peaks, it would be fairly easy to pick out some structural features of an unknown molecule.

Lists of known peaks often include molar absorptivity values as well. That might help you to be

even surer. For example using the simple carbon-oxygen double bond of ethanal (CH3CHO)

which has two peaks in its spectrum at 180 and 290nm, data shows that the peak at 290 has a

molar absorptivity of only 15, compared with the one at 180 of 10000. If your spectrum showed

a very large peak at 180, and an extremely small one at 290, that just adds to your certainty.

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3.6 Brief Introduction to A Spectrophotometer

Radiation or light source Monochromator Sample cell Detector

Meter

Figure 4: A schematic representation of a spectrophotometer

The schematic diagram above gives brief description of the basic components of a

spectrophotometer.

Radiation or light source: a light source produces a polychromatic beam of light. The UV-

Visible spectrophotometer uses deuterium, tungsten halogen or xenon lamps or LED as the light

source. Deuterium arc lamps measure in the ultraviolet (UV) region 190 370 nm. Tungsten

halogen also known as Gquartz iodine lamp measure in the visible region from 320 1100nm.

Xenon lamp measure in both the UV and visible regions of the electromagnetic spectrum from

190- 1100nm.

Monochromator: is a device which is used to resolve polychromatic radiation into its individual

wavelength and isolates these wavelengths into a very narrow band. This device produces light

radiation of only particular (single) wavelength. Here, the monochromator selects a particular

wavelength for the incident beam of light (I0).

Sample cell or cuvette: is the container that holds the sample to be analysed.

Detector: measures the intensity of the transmitted light (I).

Meter: The readout is supplied by a meter.

3.7 Brief description of how to use a UV Visible spectrophotometer

The sample is placed in the cuvette and is then irradiated with an incident beam of light (I0) of a

specific wavelength. A detector then measures the amount of light that is transmitted through the

sample (I). The signal from the detector drives a meter that can be calibrated to read

transmittance or absorbance.

To use the UV- Visible spectrophotometer, the calibration procedure entails setting 0

Absorbance at a given wavelength with a cuvette containing a reference or blank solution.

Typically, the blank solution is just the solvent, i.e., the cuvette is filled with only the solvent

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used to dissolve the sample to be analyzed, and inserted into the cell holder of the

spectrophotometer. 0 (zero) absorbance is pressed on the spectrophotometer to set the blank at 0.

After this, the blank is removed, and an identical cuvette containing the solution of interest is

then inserted into the spectrometer, and the absorbance is read from a meter on the instrument.

Both the calibration and the reading must be done at the same wavelength. The reading for the

solution then represents the absorbance at the chosen wavelength due to the component of

interest. The calibration has accounted for any absorption (or reflection or scattering) of light by

the cuvette and other species in the reference solution.

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EXPERIMENT 1 Determination of Absorption Curve and Concentration of Potassium

Nitrate

Purpose: To prepare absorption curve for potassium nitrate and use it to determine the

concentration of potassium nitrate

Discussion: Potassium nitrate is an example of an inorganic compound which absorbs mainly

in the ultraviolet. The absorbance and percentage transmittance of an approximately 0.1m

potassium nitrate solution are measured over the wavelength range 240-360 nm at 5 nm intervals

and at smaller intervals in the vicinity of the maxima or minima. Manual spectrophotometers are

calibrated to read both absorbance and percentage transmittance on the dial settings, whereas the

automatic recording double-beam spectrophotometers usually use chart paper printed with both

scales. The linear conversion chart is useful for visualizing the relationship between these two

quantities. The three normal means of presenting the spectrophotometric data are described

below; by far the most common procedure is to plot absorbance against wavelength (measured in

nanometres). The wavelength corresponding to the absorbance maximum (or transmittance

minimum) is read from the plot and used to prepare the calibration curve. This point is chosen for

two reasons: (1) it is the region in which the greatest difference in absorbance between any two

different concentrations will be obtained, thus giving the maximum sensitivity for concentration

studies, and (2) as it is a turning point on the curve it gives the least alteration in absorbance

value for any slight variation in wavelength. No general rule can be given concerning the

strength of the solution to be prepared, as this will depend upon the spectrophotometer used for

the study. Usually a 0.01-0.001M solution is sufficiently concentrated for the highest

absorbances, and other concentrations are prepared by dilution. The concentrations should be

selected such that the absorbance lies between about 0.3 and 1.5. For the determination of the

concentration of a substance, select the wavelength of maximum absorption for the compound

(e.g 302.5 to 305 nm for potassium nitrate) and construct a calibration curve by measuring the

absorbances of four or five concentrations of the substance (e.g 2, 4, 6, 8, and 10gl-1

KNO3) at

the selected wavelength. Plot absorbance against concentration. If the compound obeys Beers

law, the result will be a linear calibration curve passing through the origin. If the absorbance of

the unknown solution is measured, the concentration can be obtained from the calibration curve.

If it is known that the compound obeys Beers law, the molar absorption coefficient e can be

determined from one measurement of the absorbance of a standard solution. The unknown

concentration is then calculated using the value of the constant e and the measured value of the

absorbance under the same conditions.

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Equipment / Materials:

Potassium nitrate

Beakers

Stirring rod

Dessicator

Distilled water

UV-Visible

spectrophotometer

Experimental Procedure

Dry some pure potassium nitrate at 1100C for 2-3h and cool in a desiccator. Prepare an aqueous

solution containing 10.000gl-1

. With the aid of a spectrophotometer and matched 1cm rectangular

cells, measure the absorbance and the percentage transmittance over a series of wavelengths

covering the range 240-350 nm. Plot the data in three different ways:

(a) absorbance against wavelength

(b) percentage transmittance against wavelength

(c) log (molar absorptivity, or extinction coefficient () against wavelength.

From the curves, evaluate the wavelength of maximum absorption (or minimum transmission).

Use this value of the wavelength to determine the absorbance of solutions of potassium nitrate

containing 2.000, 4.000, 6.000 and 8.000gl-1

KNO3. Run a blank on the two cells, filling them

both with distilled water; if the cells are correctly matched, no difference in absorbance should be

discernible. Plot absorbance against concentration for each cell. Determine the absorbance of an

unknown solution of potassium nitrate and read the concentration from the calibration.

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EXPERIMENT 2 Spectrophotometric Determination of Iron

Purpose: To prepare absorption curve for iron and use it to determine the concentration of

Iron

Discussion: A complex of iron (II) is formed with 1,10-phenanthroline [Fe(C12HsN2)32+

], and

the absorbance of this coloured solution is measured with a spectrophotometer. The spectrum is

plotted to determine the absorption maximum. Hydroxylamine (as the hydrochloride salt to

increase solubility) is added to reduce any Fe3+

to Fe2+

and to maintain it in that state.

Equation. 4Fe3+

+ 2NH2OH 4Fe2+

+ N2O + 4H+ + H2O


1 litre volumetric flask Ferrous ammonium sulphate UV Visible spectrophotometer

Beakers Conc sulphuric acid Weighing balance

Stirring rod Hydroxylammonium chloride

Pipets Sodium acetate

Distilled water 1, 10 phenanthroline monohydrate

Experimental Procedure:

Solution Preparation

1. Standard iron (II) solution. Prepare a standard iron solution by weighing 0.0702 g of

ferrous ammonium sulphate, Fe(NH4)2(SO4)2.6H2O. Quantitatively transfer the weighed

sample to a one-litre volumetric flask and add sufficient water to dissolve the salt. Add

2.5 ml of concentrated sulphuric acid, dilute exactly to the mark with distilled water, and

mix thoroughly. This solution contains 10.0 mg of iron per litre (10 ppm); if the amount

weighed is different than that above, calculate the concentration.

2. 1, 10-phenanthroline solution. Dissolve 100mg of 1, 10-phenanthroline monohydrate in

100 ml of water

3. Hydroxylammonium chloride solution. Dissolve 10 g of hydroxylammonium chloride in

100 ml of water.

4. Sodium acetate solution. Dissolve 10 g of sodium acetate in 100 ml of water.

Into a series of 100 ml volumetric flasks, add with pipettes 1.00, 5.00, 10.00, and 25.00ml

17

of the standard iron solution. Into another 100 ml volumetric flask, place 50 ml of distilled water

for a blank. The unknown sample will be furnished in another 100-ml volumetric flask. To each

of the flasks (including the unknown) add 1.0ml of the hydroxylammonium chloride solution and

5.0 ml of the 1, 10-phenanthroline solution. Buffer each solution by the addition of 8.0 ml of the

sodium acetate solution to produce the red color of ferrous 1,10-phenanthroline. [The iron (ii)-

phenanthroline complex forms at pH 2 to 9. The sodium acetate neutralizes the acid present and

adjusts the pH to a value at which the complex forms.] Allow at least 15 minutes after adding the

reagents before making absorbance measurements so that the colour of the complex can fully

develop. Once developed, the colour is stable for days. Dilute each solution to exactly 100 ml.

The standards will correspond to 0.1, 0.5; 1, and 2.5 ppm of iron, respectively. Obtain the

absorption spectrum of the 2.5-ppm solution by measuring the absorbance from about 400 nm to

700 nm (or the range of your instrument). Take reading at 25-nm intervals except near the

vicinity of the absorption maximum, where you should take readings at 5 or 10 nm intervals.

Follow your instructors directions for the operation of your spectrophotometer. The blank

solution should be used as the reference solution. Plot the absorbance against the wavelength and

select the wavelength of maximum absorption. Calculate the molar absorptivity of the iron (ii)-

phenanthroline complex at the absorption maximum.

Prepare a calibration curve by measuring the absorbance of each of the standard solutions at the

wavelength of maximum absorbance. Measure the unknown in the same way. Plot the

absorbance of the standards against concentration in ppm. From this plot and the absorbance of

the unknown, determine the final concentration of iron in your unknown solution. Report the

number of micrograms of iron in your unknown along with the molar absorptivity and the

spectrum of the iron (ii)-phenanthroline complex.

18

EXPERIMENT 3 Spectrophotometric Determination of Aspirin (2-(acetylbenzoic acid)

by Iron (III)

Purpose: To determine the amount of aspirin in a commercial aspirin product.

Discussion: A coloured complex is formed between aspirin (2-(acetylbenzoic acid) and the iron

(III) ion. The intensity of the colour is directly related to the concentration of aspirin present;

therefore, spectrophotometric analysis can be used. A series of solutions with different aspirin

concentrations will be prepared and complexed. The absorbance of each solution will be

measured and a calibration curve will be constructed. Using the standard curve, the amount of

aspirin in a commercial aspirin product can be determined. The complex is formed by reacting

the aspirin with sodium hydroxide to form the salicylate dianion.


125 ml Erlenmeyer flasks Acetylsalicylic acid UV Visible spectrophotometer

2 cuvettes 50 ml volumetric flask Commercial aspirin

10 ml graduated cylinder 1 M NaOH

5 ml pipet Analytical balance

250 ml volumetric flask 0 02M Iron (III) buffer

Be careful while boiling the sodium hydroxide solution. NaOH solutions are dangerous,

especially when hot.


Part I: Making Standards. 1. Weigh 400 mg of acetylsalicylic acid (aspirin) in a 125 mL Erlenmeyer flask. Add 10 mL of a

1 M NaOH solution to the flask, and heat until the contents begin to boil.

2. Quantitatively transfer the solution to a 250 mL volumetric flask, and dilute with distilled

water to the mark.

3. Pipette a 2.5 mL sample of this aspirin standard solution to a 50 mL volumetric flask. Dilute to

the mark with a 0.02 M iron (III) solution. Label this solution "A," and place it in a 125 mL

Erlenmeyer flask.

4. Prepare similar solutions with 2.0, 1.5, 1.0, and 0.5 mL portions of the aspirin standard. Label

these "B, C, D, and E."

19

Part II: Making an unknown from a tablet.

1. Place one aspirin tablet in a 125 mL Erlenmeyer flask. Add 10 mL of a 1 M NaOH solution to

the flask, and heat until the contents begin to boil.

2. Quantitatively transfer the solution to a 250 mL volumetric flask, and dilute with distilled

water to the mark.

3. Pipette a 2.5 mL sample of this aspirin tablet solution to a 50 mL volumetric flask. Dilute to

the mark with a 0.02 M iron (III) solution. Label this solution "unknown," and place it in a 125

mL Erlenmeyer flask.

Part III: Testing the Solutions.

Turn on the spectrophotometer. Press the A/T/C button on the Spectrophotometer to select

absorbance. Adjust the wavelength to 530 nm by pressing the nm arrow up or down.

Insert the blank (0ppm cuvette of iron buffer) into the cell holder and close the door. Position

the cell so that the light passes through clear walls. *Remember to wipe off the cuvette with a

tissue paper before inserting it into the instrument.

Press 0 ABS/100% T to set the blank to 0 (zero) absorbance. Record the absorbance of the 0ppm

solution. Obtain absorbance readings for each of the other standard solutions. Record the results

on the data sheet. Obtain an absorbance reading for the unknown sample(s). Make a graph of

concentration (x-axis) vs. absorbance (y-axis). From the standard curve, determine the

concentration of aspirin in the unknown

20

4.0 CONCLUSION

UV-Visible spectroscopy refers to absorption spectroscopy. It involves the absorption of

radiation at the near ultraviolet region 200 nm to 780 nm. It is routinely used in analytical

chemistry for the quantitative and qualitative determination of different analytes, such as highly

conjugated organic compounds, transition metal ions and biological macromolecules.

5.0 SUMMARY

UV-Visible spectroscopy is an analytical technique which involves the absorption of

electromagnetic radiation at the near ultraviolet (200 nm) region and the visible region. The

ultraviolet and visible absorption is based on molecules containing pi- electrons or non bonding

electrons (n- electrons) which can absorb the energy in the form of ultraviolet or visible light to

excite these electrons to higher anti-bonding molecular orbitals. UV-Visible spectroscopy is

particularly important for the qualitative and quantitative determination of many organic

compounds especially those with a high degree of conjugation and transition metal ions. The

basic principle of quantitative determination by UV-Visible spectroscopy lies in comparing the

extent of absorption of a sample solution with that of a set of standards under radiation of a

selected wavelength through the application of Beer- Lambert law.


1. What type of electrons in a molecule is generally involved in the absorption of UV or visible

radiation.

2. Describe how the wavelength of maximum absorption can be determined.

3. What are the most frequent electronic transitions during absorption of electromagnetic

radiation which results in more intense absorption?

4. How would you obtain a calibration curve?


1. Mendham, J., Denney, R.C., Barnes, J.D., and Thomas, M.J.K., (2008) ,Vogels Textbook of



2. Gary, D.C., (1980), Analytical Chemistry, 3rd


21

UNIT 2 - COLORIMETRY

1.0 INTRODUCTION

2.0 OBJECTIVES

3.0 MAIN CONTENT

3.1 Principles of colorimetry

3.2 Colorimetric determinations

3.3 Mode of operation of a colorimeter

3.4 Experimental

4.0 CONCLUSION

5.0 SUMMARY



22

1.0 INTRODUCTION

Colorimetry is an analytical technique (spectroscopic method) used to determine the

concentrations of coloured substances in solution. It relies on the fact that a coloured

substance absorbs light (at the visible region) of a colour complementary to its own and

the amount of light it absorbs (absorbance) is proportional to its concentration .

Colorimetric determinations are carried out by the use of an instrument called colorimeter.

2.0 OBJECTIVES


Define and explain colorimetric analysis

Describe or state the principle of colorimetry

Explain the mode of operation of a colorimeter

Explain how the concentration of a coloured compound can be determined using a colorimeter

Carry out colorimetric determinations

3.0 MAIN CONTENT

3.1 Principle of colorimetry

Colorimetric analysis is based on Beer- Lambert law. The concentration of the coloured

compound is related to the amount of visible radiation absorbed by the coloured compound. The

coloured compound absorbs white light at the visible region of a colour complementary to its

own. Remember that white light is made up of different colours; red, orange, green, yellow, blue,

indigo and violet. These colours occur at particular wavelengths, e.g. blue occurs at the

wavelength 435- 480 nm, so, in colorimetry the complimentary colour absorbed by a coloured

solution occurs at particular wavelength.

The colour of the compound is usually due to the formation of a coloured compound by the

addition of an appropriate reagent or it may be inherent in the desired constituent itself. The

basic principle of most colorimetric measurements consists in comparing, under well defined

conditions, the colour produced by the substance in the unknown concentration with the same

colour produced by a known amount (standard solution) of the same substance being determined.

3.2 Colorimetric determinations

The determination is similar to spectrophotometric determination. It will normally require

1. A weighed quantity of the material under investigation in an appropriate solvent

2. A standard solution of the compound being determined in the same solvent

3. The requisite reagent

23

4. Any ancillary reagents such as buffers, acids or alkalis necessary to establish the correct

conditions for formation of the required coloured product

5. Preparation of calibration curve

6. Estimation of the unknown concentration of the test solution from the calibration curve.

In view of the sensitivity of colorimetric and spectrophotometric methods, the absorbance

measurements are usually made on very dilute solutions. In order to take sufficient material for

an accurate weight to be achieved when preparing the original solution of analyte and the

corresponding standard solution, it is commonly necessary to prepare solutions which are too

concentrated for the absorbance measurements, and these must then be diluted accurately to the

appropriate strenght.

Solutions of colour producing reagents are frequently unstable and normally should not be stored

for more than a day or so.

The colorimeter consists of light source, monchromator, slit, optical cell or cuvette, photoelectric

cell and galvanometer.

3.3 Mode of operation

White light from a tungsten lamp passes through a condenser lens to give a parallel beam

which falls on the filter that is positioned to select radiation of specific wavelength to impinge

on a glass cuvette containing the solution. As the light is passing through the solution, some part

of it is absorbed by the sample component, while the part that is not absorbed is transmitted, and

detected by a photo electric cell (detector). In order to measure the absorbance of a solution, the

meter reading is first adjusted to 100% transmittance (zero absorbance) with a blank solution.

The sample is then inserted in place of the blank and the absorbance is read directly. The

concentration corresponding to the absorbance of the sample is then obtained from the standard

or calibration curve. The filter is usually a complimentary colour of the test solution.

The filter is chosen to select the band of wavelengths which are most strongly absorbed by the

coloured solution e.g. this is illustrated in the table below, by using a yellow filter to use in

measuring the concentration of a blue coloured solution like copper (II) sulphate or its

ammine/amine complex.

24

Table 1

The wavelength (nm) of the observed transmitted colour of the solution

The observed transmitted colour of the solution

The complementary colour of the solution i.e. the colour

of the filter

400-435 violet yellowish-green

435-480 blue * yellow *

480-490 greenish-blue orange

490-500 bluish-green red

500-560 green purple

560-580 yellowish-green violet

580-595 yellow blue

595-610 orange greenish-blue

610-750 red bluish-green

25

EXPERIMENT 1 Colorimetric determination of manganese in steel

Purpose: To determine the concentration of manganese in steel

Discussion: Colorimetry is particularly suited to the determination of manganese in

steel because the manganese can be converted into permanganate ions, which are

coloured. The conversion is achieved in two stages. Using nitric acid, the managanese is

first oxidised to manganese (II) ions, which are then oxidised to permanganate ions by

the more powerful oxidising agent, potassium periodate.


Measuring cylinders (50cm3

and 10 cm

3)

Dropper Potassium persulphate

Clock glass Wire cutters Standard flask (50 cm3 and 100 cm

3)

Filter funnel Propanone Potassium permanganate

Tweezers Deionised water Green filter

Wash bottle Anti bumping granules

Optical matched cuvettes

Glass beakers (50 cm3 and 250 cm

3) Analytical balance

(accurate to 0.001g)

Bunsen burner, tripod stand,wire

gauze

Steel paper chips Colorimeter Acidified 85% phosphoric acid

Acidified potassium periodate(5g

potassium periodate per 100 cm3 of 2

mol 1-1

nitric acid

CAUTION

Wear eye protection and if any chemical splashes on your skin wash it off immediately.

The acidified 0.0010 mol l1

potassium permanganate is harmful if ingested and irritates

the eyes and skin. Wear gloves.

Both 2 mol l1

nitric acid and its vapour are corrosive and toxic, causing severe burns to

the eyes, digestive and respiratory systems. Wear gloves.

85% phosphoric acid is corrosive: it burns and irritates the eyes and skin. It is a

systemic irritant if inhaled and if swallowed causes serious internal injury. Wear

gloves.

Acidified potassium periodate solution is harmful if swallowed and is an irritant to the

eyes, skin and respiratory system. It is also corrosive. Wear gloves.

Potassium persulfate is harmful if swallowed or inhaled a s a dust. It irritates the eyes,

skin and respiratory system, causing dermatitis and possible allergic reactions. Wear

gloves.

26

Propanone is volatile and highly flammable, and is harmful if swallowed. The vapour

irritates the eyes, skin and lungs, and is narcotic in high concentrations. Wear gloves.

Experimental Procedure

Part A Calibration graph

1. Rinse the burette, including the tip, with 0.0010 mol l1 acidified potassium

permanganate and fill it with the same solution.

2. Run 2 cm3 of the permanganate solution into a 50 cm3 standard flask and make

up to the graduation mark with deionised water.

3. Stopper the flask and invert it several times to ensure the contents are completely

mixed.

4. Rinse a cuvette with some of the solution and f ill it.

5. Using a colorimeter (fitted with a green filter) measure the absorbance of the

solution in the cuvette. If you have more than one green filter, choose the one that

gives maximum absorbance.

6. Repeat steps 2 to 5 with 4, 6, 8, 10, 12 and 14 cm3 of the permanganate stock

solution in the burette.

7. Use dilution formular to calculate the concentration of the 2cm3 14 cm3 of the

diluted stock solution.

8. Plot a calibration graph of absorbance against concentration of potassiu m

permanganate.

Part B Conversion of manganese to permanganate

1. Degrease a steel paper clip by swirling it with a little propanone in a beaker.

Using tweezers remove the paper clip and leave it to dry for a minute or so on a

paper towel.

2. Cut the paper clip into small pieces.

3. Weigh accurately about 0.2 g of the paper clip pieces and transfer them to a 250

cm3 glass beaker.

4. Add approximately 40 cm3 of 2 mol l1 nitric acid to the beaker and cover it with

a clock glass.

5. Heat the mixture cautiously, in a fume cupboard, until the reaction starts.

Continue heating gently to maintain the reaction, but remove the source of heat if

the reaction becomes too vigorous.

6. Once the steel has reacted, allow the solution to cool a little. Add a couple of anti-

27

bumping granules and then boil the solution until no more brown fumes are given

off.

7. Once this solution has cooled considerably no more than hand hot adds about 5

cm3 of 85% phosphoric acid, approximately 0.2 g of potassium persulfate and a

couple of fresh anti-bumping granules. Boil the mixture for about 5 minutes.

8. To this solution, add approximately 15 cm3 of acidified potassium periodate solution

plus a couple of fresh anti-bumping granules and then gently boil the mixture. The

solution will start to turn pink. Continue gently boiling until the intensity of the pink

colour remains constant. This should take about 5 minutes.

9. Allow the pink solution to cool to room temperature and then transfer it to a 100 cm3

standard flask, leaving the anti-bumping granules in the beaker.

10. Rinse the beaker several times with a little deionised water and add the rinsings (but

not the anti-bumping granules) to the flask.

11. Make up the solution to the graduation mark with deionised water.

12. Stopper the flask and invert it several times to ensure the contents are completely

mixed.

13. Using a colorimeter fitted with the appropriate green filter, measure the absorbance

of the solution.

14. Use your calibration graph to convert the absorbance to permanganate

concentration and then calculate the percentage by mass of manganese in the steel

paper clip.

28

4.0 CONCLUSION

Colorimetry is a spectroscopic method of analysis which involves the measurement of the

absorption of electromagnetic radiation at the visible region by a coloured compound. The

principle is based on the application of Beer- Lambert law. The amount of visible radiation

absorbed by the coloured compound is related to the concentration of the analyte in the solution.

5.0 SUMMARY

Colorimetry is an analytical technique (spectroscopic method) used to determine the

concentrations of coloured substances in solution. It relies on the fact that a coloured

substance absorbs light at the visible region of the electromagnetic radiation , of a colour

complementary to its own and the amount of light it absorbs (absorbance) is

proportional to its concentration. Colorimetric determinations are carried out by the use of an instrument called colorimeter. The unknown concentration of the test sample solution is obtained

from the calibration curve of the standard solutions of the test solution.


1. In what region of the electromagnetic spectrum is colorimetric analysis carried out

2. Why is it necessary to prepare fresh solution of the coloured sample to be analysed

3.0 Describe the mode of operation of a colorimeter

4. Can colorimeter be used to analyse dilute coloured solutions .




Edition. Pearson Education.



3. Colorimetry, http://www.docbrown.info/page07/appendixtrans09.htm.

http://www.docbrown.info/page07/appendixtrans09.htm

29

UNIT 3 - INFRARED SPECTROSCOPY

1.0 INTRODUCTION

2.0 OBJECTIVES

3.0 MAIN CONTENT

3.1 Principle of Infrared (IR) spectroscopy

3.2 Types of vibrations

3.3 Group frequencies

3.4 Correlation of structure and frequency

3.5 IR Spectroscopy Experimental Procedures

3.6 Experiments

4.0 CONCLUSION

5.0 SUMMARY



30

1.0 INTRODUCTION

Infrared spectroscopy is a spectroscopic method of analysis used qualitatively to identify and

study chemicals and quantitavely to measure concentration. When a molecule interacts/absorbs

radiation at the infrared region, it causes the molecule to undergo vibrational transitions. Let us

pause here to get a clearer picture or understanding of what happens when a molecule absorbs

radiation.

There are three basic processes by which a molecule can absorb radiation: all involve raising the

molecule to a higher internal energy level, the increase in energy being equal to the energy of the

absorbed radiation (hv). The three types of internal energy are quantized; that is, they exist at

discrete levels. Firstly, the molecules rotates about various axes, the energy of rotation being at

definite energy levels, so the molecule may absorb radiation and be raised to a higher rotational

energy level, in a rotational transition. This type of transition (rotational transition) occurs when

molecules absorb radiation at the far infrared region and microwave region of the

electromagnetic spectrum. Secondly, the atoms or groups of atoms within a molecule vibrate

relative to each other, and the energy of this vibration occurs at definite quantized levels. The

molecule may then absorb a discrete amount of energy and be raised to a higher vibrational

energy level, in a vibrational transition. This type of transition occurs with absorption of near

infrared radiation. Third, the electrons of a molecule may be raised to a higher electron energy,

corresponding to an electronic transition. This type of transition occurs at the ultraviolet and

visible region. These transitions occur only at definite wavelengths corresponding to an energy

equal to the difference of the discrete energy levels involved in the transition.

The infrared portion of the electromagnetic spectrum is usually divided into three regions; the

near-, mid- and far- infrared. The higher-energy near-IR, approximately 140004000 cm1

(0.8

2.5 m wavelength) can excite overtone or harmonic vibrations. The mid-infrared,

approximately 4000400 cm1

(2.525 m) may be used to study the fundamental vibrations and

associated rotational-vibrational structure. The far-infrared, approximately 40010 cm1

(25

1000 m), lying adjacent to the microwave region, has low energy and may be used for

rotational spectroscopy. Most of the analytical applications are confined to the middle IR region

because absorption of organic molecules are high in this region.

The method or technique of infrared spectroscopy uses an instrument called an infrared

spectrometer to produce an infrared spectrum. A basic IR spectrum is essentially a graph of

infrared light absorbance (or transmittance) on the vertical axis vs. frequency or wavelength on

the horizontal axis. Typical units of frequency used in IR spectra are reciprocal centimeters

(sometimes called wave numbers), abbreviated as cm1

. Units of IR wavelength are commonly

given in microns, abbreviated as m, which are related to wave numbers in a reciprocal way.

http://en.wikipedia.org/wiki/Overtoneshttp://en.wikipedia.org/wiki/Harmonic_(mathematics)http://en.wikipedia.org/wiki/Rotational-vibrational_spectroscopyhttp://en.wikipedia.org/wiki/Microwavehttp://en.wikipedia.org/wiki/Rotational_spectroscopyhttp://en.wikipedia.org/wiki/Absorbancehttp://en.wikipedia.org/wiki/Transmittancehttp://en.wikipedia.org/wiki/Units_of_measurementhttp://en.wikipedia.org/wiki/Reciprocal_centimetershttp://en.wikipedia.org/wiki/Wave_numberhttp://en.wikipedia.org/wiki/Micronhttp://en.wikipedia.org/wiki/Multiplicative_inverse

31

2.0 OBJECTIVES


Explain the principle of infrared spectroscopy

Give account of the kinds of molecules that absorb infrared radiation

State the types of vibration

Carry out IR spectroscopy practical

3.0 MAIN CONTENT

3.1 Principle of Infrared Spectroscopy

When molecules absorb radiation at the infrared region, the energy of the wavelength absorbed

causes a vibrational transition if the energy absorbed is equal to the quantized jump in the

internal energy i.e. the energy difference between the vibrational energy levels of the atoms or

group of atoms within the molecule. The vibrational energy of the atoms or groups of atoms

within the molecule is raised to a higher level. The energy absorbed appears as absorption peaks

at the wavelength it corresponds to on the IR spectrum.

Not all molecules can absorb in the infrared region. For absorption to occur there must be change

in the dipole moment (polarity) of the molecule. A diatomic molecule must have a permanent

dipole (polar covalent bond in which a shared pair of electrons is shared unequally) in order to

absorb.

Different functional groups absorb characteristics frequencies of IR radiation. Hence gives the

characteristic peak value. Therefore IR spectrum of a chemical substance is a finger print of a

molecule for its identification.

3.2 Types of Vibrations

In an organic molecule there are two major types of fundamental vibrations. These are stretching

and bending vibrations. The stretching vibration could either be symmetrical or asymmetrical.

Bending vibration is of four different types namely scissoring, rocking, wagging and twisting.

The energy required to bend a bond is not great and falls within the range of 400 1300 cm-1

.

This region is called the finger print region because absorption in this region is very dependent

on the molecular environment. Thus, this region is used to establish the identity of the chemical

compounds. The energy required to stretch a bond is a little bit higher. This falls within the

region 0f 1300 4000 cm-1

. Absorption in this region is caused by functional groups and is

independent of other parts of the molecule and is used to detect the functional groups in

molecules.

3.3 Group frequencies

Group frequencies are the absorption bands or signals that occurs at certain frequencies due to

stretching or bending vibration within a molecule. For example, the bands at 3300 cm-1 and

1050 cm-1 are characteristics of the OH group in alcohols.

32

3.4 Correlation of structure and frequency

Many thousands of infrared spectra have been recorded and from these, it has been possible to

empirically tabulate correlations between absorption frequencies and types of bonds or chemical

groups. Table 1 summaries some of the correlations for various types of vibrations.

Table 2: Group Frequencies

Vibration Type of molecule Group frequencies (cm-1)

C H stretch Alkanes, alcohols 2800 - 3000

C H stretch Aldehydes 2700 - 2900

C H stretch Alkenes 3010 - 3095

O H stretch Alcohols, phenols 3200 - 3600

O H stretch Acids 2500 - 3000

O H bend Alcohol, phenol 1260 - 1410

N H stretch Amines 3300 - 3500

C = C stretch Alkenes 1620 - 1680

C = O stretch Aldehydes 1720 - 1740

C = C stretch Alkynes 2100 - 2140

C = Nstretch Nitriles 2000 - 2500

C = Ostretch Ketones 1705 - 1725

3.5 IR Spectroscopy Experimental Procedures

The infrared spectrum of a sample is recorded by passing a beam of infrared light through the

sample. When the frequency of the IR is the same as the vibrational frequency of a bond,

absorption occurs. The energy absorbed appears as absorption peaks at the wavelenght it

corresponds to on the IR spectrum. A basic IR spectrum is essentially a graph of infrared light

absorbance (or transmittance) on the vertical axis vs. frequency or wavelength on the horizontal

axis. Well resolved and sharp peak/peaks at wavelength/wavelenghts it corresponds to, is/are

matched with the wavelenght range on the group frequency table to identify the functional

group/groups present.

Analysis of a sample by IR spectroscopy involves:

Sample preparation

Gaseous samples require a sample cell with a long pathlength to compensate for the diluteness.

The pathlength of the sample cell depends on the concentration of the compound of interest. A

simple glass tube with length of 5 to 10 cm equipped with infrared-transparent windows at the

both ends of the tube can be used for concentrations down to several hundred ppm. Sample gas

concentrations well below ppm can be measured with a White's cell in which the infrared light is

http://en.wikipedia.org/wiki/Absorbancehttp://en.wikipedia.org/wiki/Transmittancehttp://en.wikipedia.org/wiki/Pathlengthhttp://en.wikipedia.org/wiki/White_cell_(spectroscopy)

33

guided with mirrors to travel through the gas. White's cells are available with optical pathlength

starting from 0.5 m up to hundred meters.

Liquid samples can be sandwiched between two plates of a salt (commonly sodium chloride, or

common salt, although a number of other salts such as potassium bromide or calcium fluoride are

also used).The plates are transparent (do not absorb in the IR region) to the infrared light and do

not introduce any lines onto the spectra.

Solid samples can be prepared in a variety of ways. One common method is to crush the sample

with an oily mulling agent (usually Nujol) in a marble or agate mortar, with a pestle. A thin film

of the mull is smeared onto salt plates and measured. The second method is to grind a quantity of

the sample with a specially purified salt (usually potassium bromide) finely (to remove scattering

effects from large crystals). This powder mixture is then pressed in a mechanical press to form a

translucent pellet through which the beam of the spectrometer can pass. A third technique is the

"cast film" technique, which is used mainly for polymeric materials. The sample is first dissolved

in a suitable, non hygroscopic solvent. A drop of this solution is deposited on surface of KBr or

NaCl cell. The solution is then evaporated to dryness and the film formed on the cell is analysed

directly.

Simple spectra are obtained from samples with few IR active bonds and high levels of purity.

More complex molecular structures lead to more absorption bands and more complex spectra.

Comparing to a reference

To take the infrared spectrum of a sample, it is necessary to measure both the sample and a

"reference" (or "control"). This is because each measurement is affected by not only the light-

absorption properties of the sample, but also the properties of the instrument (for example, what

light source is used, what infrared detector is used, etc.). The reference measurement makes it

possible to eliminate the instrument influence. Mathematically, the sample transmission

spectrum is divided by the reference transmission spectrum.

The appropriate "reference" depends on the measurement and its goal. The simplest reference

measurement is to simply remove the sample (replacing it by air). However, sometimes a

different reference is more useful. For example, if the sample is a dilute solute dissolved in water

in a beaker, then a good reference measurement might be to measure pure water in the same

beaker. Then the reference measurement would cancel out not only all the instrumental

properties (like what light source is used), but also the light-absorbing and light-reflecting

properties of the water and beaker, and the final result would just show the properties of the

solute (at least approximately).

A common way to compare to a reference is sequentially: first measure the reference, then

replace the reference by the sample and measure the sample.

http://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Potassium_bromidehttp://en.wikipedia.org/wiki/Calcium_fluoridehttp://en.wikipedia.org/wiki/Nujolhttp://en.wikipedia.org/wiki/Marblehttp://en.wikipedia.org/wiki/Agatehttp://en.wikipedia.org/wiki/Mortar_and_pestlehttp://en.wikipedia.org/wiki/Potassium_bromidehttp://en.wikipedia.org/wiki/Machine_presshttp://en.wikipedia.org/wiki/Infrared_detector

34

3.6 Measuring IR absorption bands

As with electronic (uv-visible) spectra, the use of infrared spectra for quantitative determinations

depends upon measuring the intensity of either the transmission or absorption of the infrared

radiation at a specific wavelength, usually the maximum of a strong, sharp, narrow, well-

resolved absorption band. Most organic compounds will possess several peaks in their spectra

which satisfy these criteria and which can be used so long as there is no substantial overlap with

the absorption peaks from other substances in the sample matrix.

The background to any spectrum does not normally correspond to a 100% transmittance at all

wavelengths, so measurements are best made by what is known as the baseline method. This

involves selecting an absorption peak to which a tangential line can be drawn, as shown in

Figure. This is then used to establish a value for I0 by measuring vertically from the tangent

through the peak to the wavenumber scale. Similarly, a value for I is obtained by measuring the

corresponding distance from the absorption peak maximum. So, for any peak, the absorbance

will not be the value corresponding to the height of the absorption, measured from the horizontal

axis of the chart paper; instead it will be the value of Acalc obtained from the equation

Acalc =

=

Where I0 and I are values measured using the tangential baseline.

This procedure has the great advantage that some potential sources of error are eliminated. The

measurements do not depend upon accurate wavelength positions as they are made with respect

to the spectrum itself, and any cell errors are avoided by using the same cell of fixed path length.

Measuring Acalc eliminates any variation in the source intensity, the instrument optics or the

sensitivity.

3.7 Beers law: Quantitative IR spectra

Infrared spectra are recorded using either or both absorbance (A) and percentage transmittance

(T) just as they are in visible ultraviolet electronic spectra, and Beers law,

A = cl =

=

applies equally to infrared spectra as it does to electronic spectra.

Use of a calibration graph

A calibration curve overcomes any problems created due to non-linear absorbance or

concentration features, and it means that any unknown concentration run under the same

conditions as the series of standards can be determined from the graph. The procedure requires

that all standards and samples are measured in the same cell of fixed path length, although the

dimensions of the cells and the molar absorptivity for the chosen absorption band are not needed;

they are constant for all the measurements.

35

EXPERIMENT 1 The Identification of Functional Groups of an Unknown Substance

Purpose: This experiment involves the instrumental technique to obtain the spectrum as well as

the analysis of the spectral data to determine the functional groups present in the unknown

sample.

Discussion:

A critical part of the infrared experiment is getting the infrared radiation to interact with the

sample without losing a significant part of the infrared radiation from non-sample interactions

(mirrors which absorb light rather than reflect, scratches on any optical surface which reflect

light in the wrong direction, sample surfaces which reflect light, etc.). Classically, liquids are

analyzed either neat (suspended between two sodium chloride plates (sodium chloride does not

absorb infrared radiation in the spectral range of concern)) or in solution in a solvent such as

carbon tetrachloride which does not absorb too much infrared radiation in the spectral range of

concern. Solids, with problems associated with crystal surfaces reflecting away most of the

radiation, are analyzed either in solution (at least the few solids which dissolved in suitable IR

solvents) or more commonly as a mull, a KBr pellet, or a melt (if the melting point was low

enough). Such sampling techniques are necessary to provide adequate sample for the classical

dispersive optical-null double-beam prism and (later) grating spectrometers. However, the ready

availability of low cost powerful laboratory computing has allowed the routine utilization of

more sensitive nondispersive Fourier Transform Infrared Spectrometers (FT-IR). FT-IR allows

for the collection of all the spectral data in seconds compared to 3-10 minutes for the classical

grating spectometer. However, FT-IR requires powerful computing to mathematically analyze

the collected data. Computers also allow for the collection of many spectra in a short time and

the averaging of the spectra to eliminate most random noise.


NOTE: Suggestions (detailed instructions for the operation of our spectrometer will be given by

your instructor in lab).

Normally you will not have to run a "background". (A background measures the amount of

energy that actually gets to the detector without any sample in the sampling device. The energy

reaching the detector is not constant at all wavelengths due to: absorption by spectrometer

mirrors and windows; scattering by flaws, scratches, and dirt; absorption by condensed

compounds and the atmosphers; variable output by the infrared source; etc. The computer in the

spectrometer subtracts the background from your sample data to produce the spectrum.)

However, if you are the first person to use the spectrometer for the day, or if the spectrometer has

been run for quite awhile, or if your spectrum is problematic, you should run a background.

Obtaining a spectrum of a liquid using salt plates.

a. Obtain the salt plates, holder, and O-rings from the desiccator.

36

b. Clean the salt plates by wiping with a Kim-wipe moistened with absolute alcohol. Be

extremely careful not to touch the surface of the salt plates with your fingers. (Your

fingers are wet enough to dissolve the salt!)

c. Place an O-ring on the salt plate holder.

d. Place a clean salt plate on the O-ring. Again, do not touch the flat surface of the salt

plate!!

e. Place a drop or two of a dry liquid sample on the salt plate.

f. Place the other clean salt plate on top of the sample. Be sure there are no air bubbles in

the sample.

g. Place the other O-ring on the upper salt plate.

h. Place the other metal holder on the O-ring.

i. Lightly tighten the four knurled nuts. If you tighten too hard you may crack the salt

plates or squeeze all of your sample out from between the salt plates.

j. Place the holder in the spectrometer. Obtain your spectrum. Print a hard copy and save

your spectrum on the network.

k. Disassemble the sample holder and clean the salt plates again by wiping with a Kim-

wipe moistened with absolute alcohol.

l. Return the clean salt plates in their container to the desiccator. Also, return the salt

plate holder and O-rings to the deciccator.

From the IR spectrum, determine the functional group of your unknown.

37

EXPERIMENT 2 Determination of Concentration OF Cyclohexane Using IR

Spectroscopy

Purpose: To determine the concentration of cyclohexane using IR spectroscopy


Run infrared spectra for pure cyclohexane and pure nitromethane. From the spectra select a

cyclohexane absorption which is not affected by, or overlapping with, those of the nitromethane.

Prepare a series of solutions of known concentrations of cyclohexane in nitromethane covering

the range from 0% to 20%(w/v). Using a cell of fixed path length 0.1 mm, measure the

absorbance for the solutions at the chosen peak absorption using the baseline method and plot the

calibration graph. Use this graph to determine the unknown concentration of cyclohexane in the

sample.

38

4.0 CONCLUSION

Infrared spectroscopy is a useful analytical technique used in the identification of functional

groups in a molecule. It is also useful for quantitative analysis of complex mixtures of similar

compounds because some absorption peaks for each compound will occur at a definite and

selective wavelength, with intensities proportional to the concentration of absorbing species.

Infrared spectroscopy exploits the fact that molecules absorb specific frequencies that are

characteteristic of their structure. These absorptions are resonant frequencies i.e the frequency of

the absorbed radiation matches or is equal to the transition energy of the bonds or group that

vibrates.

5.0 SUMMARY

IR spectroscopy is concerned with the study of absorption of infrared radiation, which causes

vibrational transition in the molecule. Hence, IR spectroscopy is also known as vibrational

spectroscopy. IR spectra are mainly used in structure elucidation to determine the functional

groups. The IR region is subdivided into three; the near, middle and far infrared regions. Most of

the analytical applications are confined to the middle IR region because absorption of organic

molecules are high in this region. When the applied energy in the form of infrared radiation is

equal to the vibrational transition energy of the atoms in the molecule, absorption of IR radiation

takes place and a peak is observed. Different functional groups absorb characteristic frequencies

of IR radiation. Hence gives the characteristic peak value. Therefore, IR spectrum of a chemical

substance is a finger print of a molecule for its identification. The criteria for a compound to

absorb IR radiation are; Correct frequency/wavelength of radiation that has the right amount of

energy that matches the vibrational transition energy of the atoms or group of atoms involved,

and change in dipole moment of the atoms or group of atoms involved.


1. State two criteria necessary for a molecule to absorb infrared radiation

2. Explain briefly infrared spectroscopy

3. What are group frequencies?

4. Explain how you can obtain the functional group present in a molecule

5. State the uses of infrared spectroscopy.

39







3. Infrared Spectroscopy, http:people.uwplatt.edu/~sundin/351/351h-ir.htm

40

UNIT 4 ATOMIC ABSORPTION SPECTROSCOPY

1.0 INTRODUCTION

2.0 OBJECTIVES

3.0 MAIN CONTENT

3.1 Principle of Atomic Absorption Spectroscopy

3.2 Experimental Preliminaries

3.2.1 Preparation of calibration curve

3.2.2 Preparation of sample solutions

3.2.3 Preparation of standard solutions

3.3 Uses of Atomic Absorption Spectroscopy

3.4 Experimental

4.0 CONCLUSION

5.0 SUMMARY



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NATIONAL OPEN UNIVERSITY OF NIGERIA SCHOOL OF … 391.pdf · i national open university of nigeria school of science and technology course code: chm 391 course title: practical chemistry